Lithium tetrafluoroborate

From Wikipedia, the free encyclopedia
Jump to: navigation, search
Lithium tetrafluoroborate
Li+.svg
Tetrafluoroborat-Ion.svg
Names
IUPAC name
Lithium tetrafluoroborate
Other names
Borate(1-), tetrafluoro-, lithium
Identifiers
14283-07-9 YesY
ChemSpider 3504162 YesY
Jmol-3D images Image
PubChem 4298216
Properties
LiBF4
Molar mass 93.746 g/mol
Appearance White/grey crystalline solid
Odor odorless
Density 0.852 g/cm3 solid
Melting point 296.5 °C (565.7 °F; 569.6 K)
Boiling point decomposes
Very soluble[1]
Hazards
SDS External MSDS
Main hazards Harmful, causes burns,
hygroscopic.
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentine Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium Special hazards (white): no codeNFPA 704 four-colored diamond
Related compounds
Other anions
Tetrafluoroborate,
Related compounds
Nitrosyl tetrafluoroborate
Except where noted otherwise, data is given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
 YesY verify (what isYesY/N?)
Infobox references

Lithium tetrafluoroborate is an inorganic compound with the formula LiBF4. It is a white crystalline powder. It has been extensively tested for use in commercial secondary batteries, an application that exploits its high solubility in nonpolar solvents.[2]

Applications[edit]

Although BF4 has high ionic mobility, solutions of its Li+ salt are less conductive than other less associated salts.[2] As an electrolyte in Lithium-ion batteries, LiBF4 offers some advantages relative to the more common LiPF6. It exhibits greater thermal stability[3] and moisture tolerance.[4] For example LiBF4 can tolerate a moisture content up to 620 ppm at room temperature whereas LiPF6 readily hydrolyzes into toxic POF3 and HF gases, often destroying the battery's electrode materials. Disadvantages of the electrolyte include a relatively low conductivity and difficulties forming a stable solid electrolyte interface with graphite electrodes.

Thermal stability[edit]

Because LiBF4 and other alkali-metal salts thermally decompose to evolve boron trifluoride, the salt is commonly used as a convenient source of the chemical at the laboratory scale:[5]

LiBF4LiF + BF3

Production[edit]

LiBF4 is a byproduct in the industrial synthesis of diborane:[5][6]

8 BF3 + 6 LiHB2H6 + 6 LiBF4

LiBF4 can also be synthesized from LiF and BF3 in an appropriate solvent that is resistant to fluorination by BF3 (e.g. HF, BrF3, or liquified SO2):[5]

LiF + BF3 → LiBF4

References[edit]

  1. ^ GFS-CHEMICALS
  2. ^ a b Xu, Kang. "Nonaqueous Liquid Electrolytes for Lithium-Based Rechargeable Batteries."Chemical Reviews 2004, volume 104, pp. 4303-418. doi:10.1021/cr030203g
  3. ^ S. Zhang, K. Xu, T. Jow (2003). "Low-temperature performance of Li-ion cells with a LiBF4-based electrolyte" (PDF). Journal of Solid State Electrochemistry 7 (3): 147–151. doi:10.1007/s10008-002-0300-9. Retrieved 16 February 2014. 
  4. ^ S. S. Zhang;z K. Xu; and T. R. Jow (2002). "Study of LiBF4 as an Electrolyte Salt for a Li-Ion Battery" (PDF). Journal of The Electrochemical Society 149 (5): A586–A590. doi:10.1149/1.1466857. Retrieved 16 February 2014. 
  5. ^ a b c Robert, Brotherton; Joseph, Weber; Clarence, Guibert; and John, Little (2000). "Boron Compounds". Ullmann's Encyclopedia of Industrial Chemistry: pg. 10. doi:10.1002/14356007.a04_309. 
  6. ^ Brauer, Georg (1963). Handbook of Preparative Inorganic Chemistry Vol. 1, 2nd Ed. Newyork: Academic Press. p. 773. ISBN 978-0121266011.